Atomic Orbits Home Atomic Orbits The nature of atoms and the manner in which electrons interact and move about their nucleus has puzzled scientists for a long time. What ought to be a simple orbit of electrons around a nucleus has been complicated by the fact that electrons appear to group themselves at certain fixed altitudes, while occasionally radiating light at particular frequencies. Various models of the atom have been forwarded in an attempt to explain the anomalies. They range from simple planetary models, to standing particle-waves, to probability distribution functions whose precise physical nature confounds the imagination. In earlier chapters we studied the nature of the electromagnetic force in much detail and discovered that this force has unexpected properties when charged particles are in motion. These include properties such as a sideways force which moves particles non-directly. Could the proposed modifications to Coulombs force law provide any insight to the behaviour of atoms? Here we investigate that possibility. The Rutherford model A brief history lesson is appropriate. In 1911 Ernest Rutherford studied experiments involving bombarding atoms with alpha-beams. They showed that most of an atoms mass was positively charged and packed into a small volume.

From this he developed a planetary model of the atom which proposed that electrons orbit a heavy positive nucleus; much like planets orbit the Sun. Prior to this it was believed that the nucleus was large and electrons were imbedded within: the so-named plum pudding model. Rutherfords model was intuitive. One could easily imagine a miniature solar system with the electrostatic force substituted for gravity. But the model had some problems. First, an electron orbiting a proton would radiate energy and this, according to energy conservation laws, would require it to steadily lose altitude; causing it to spiral into the nucleus. Second, electrons could orbit at any altitude and radiate any frequency. This was inconsistent with observation. The Bohr model Enter Niels Bohr, who in 1913 proposed a model of the atom based on the idea that electrons behave as waves. In this model electron waves would surround the nucleus in a standing wave formation. The wavelength of an electron was said to be based on its velocity and, seeing that velocity was also based on distance from the nucleus, it followed that electrons could only orbit at certain fixed radii.

By combining orbital mechanics, Coulombs law, and the recent work of Planck, Bohr calculated the orbital altitudes of electrons in a hydrogen atom should be: Where n is an integer number greater than zero, Rn is the nth radius, and e, k and h are the standard universal constants: electron charge, Coulomb and Planck. Bohrs model predicted a shell structure of the atom where n represents the shell number. A value of n=1 represents the shell nearest the nucleus. It is called the Bohr radius and has a value of 5.29x10-11 m. This is close to the empirical (measured) radius of a hydrogen atom: 2.5x10-11 m. Knowing the radius, he could also determine the velocity and kinetic energy of an electron at each shell. Bohr went farther to state that when an electron moves between shells it emits radiation, and the energy of this radiation will correspond to the difference between these energy levels. This energy can then be divided by Planks constant to determine frequency. It yields: Where n1 and n2 are positive integers corresponding to shells, with n2>n1.

This formula amazingly matched the observed frequencies that were emitted by excited hydrogen gas, and even more amazingly, it predicted frequencies outside the visible spectrum that had not yet been discovered. The fact that Bohrs model gets the frequencies right strongly indicates that its close to reality. Bohr was definitely on to something! Yet there are some problems with his model: According to this model, when the electrons orbit the nucleus they dont emit radiation. Realistically we should expect the same emission problem as with Rutherfords model. This lack of emission suggests the electron isnt moving. Yet the calculation for Bohrs radius is based on centrifugal force and particle-wavelength, both of which require velocity. The model says that when electrons jump between shells they emit a photon of single frequency. Single frequency means a pure sine wave. Producing pure sine electromagnetic waves requires oscillating a charged particle, e.g. an electron, in a sine wave motion  at least thats how it works for radio transmission. But an electron jumping between shells is unlikely to do so in a sine wave fashion. Instead, the radiation we would expect from a jumping movement is a step function which corresponds to multiple frequencies.

Finally (to state the obvious problem) the model doesnt explain why electrons should move as waves. The origin of emissions To gain a better understanding of the atom, a good first step may be to look at how these single frequency emissions are occurring. Bohr claims they are generated when electrons jump between shells. Yet as pointed out above, such a movement would produce only a blip in the electric field, not a sine wave. We know that producing a pure frequency in the radio portion of the electromagnetic spectrum requires oscillating electrons in a sine wave motion; e.g. within an antenna. Thus it seems reasonable to assume that producing pure frequencies in the visible portion of the spectrum requires oscillating an electron at a visible-light frequency. Are there any electrons oscillating within an atom? There sure are! How about the orbital motion of an electron around the nucleus? When viewed from the side, an electron orbiting a proton in a circle would appear as a sine wave. In fact it would appear as a sine wave from every direction. Could this orbital motion be the true origin of sinusoidal waves? We know from experiment that the radius of a hydrogen atom is approximately 2.5x10-11 m.

If an electron were to orbit at that distance, at what frequency would it do so? Taking the rule that Coulomb force should cancel centrifugal force, and the relation between frequency and velocity, namely: we get frequency as: Substituting the measured radius we determine a frequency of 2x1016 Hz, which corresponds to a wavelength of 15 nm  a wavelength in the high ultraviolet range. The observed hydrogen emission spectrum is typically in the range of 95  1000+ nm. Admittedly the calculated value of 15 nm is somewhat outside this range; yet not terribly far out. A problem may be that the measured radius was taken from a hydrogen molecule, which emits no radiation. It will be shown further on that the hydrogen which emits radiation has a different form and thus the above calculation would only be approximate in this case. Based on the above reasoning I would like to suggest that these fixed frequency emissions are not caused by electrons jumping between shells, rather, they are caused by electrons orbiting a nucleus at fixed altitudes.

Formula (1) also allows us to determine the velocity as: Using the measured radius this gives a velocity of 3.2x106 m/s, which is about 1% of light speed. This information will be relevant further on. A Simple Hydrogen Atom Following in Bohrs footsteps came Schr�dinger and Heisenberg, who postulated that electrons move according to wave functions that describe a probability distribution around the nucleus. Given that these models lay outside the classical realm, Ill not attempt to analyse their validity. Instead Ill propose a new model of the hydrogen atom. Based on the Velocity Dependant Coulombs Law (VDCL) and sideways forces, this is how an electron will orbit a proton: Here we see the electron is spiralling away from the proton. Why is this? Refer to the chapter on the sideways force . When an electron moves past the proton, force lines pass through it at an angle. This causes the electron to receive a gentle push in its direction of motion; causing it to steadily spiral away from the proton. Critics of the Rutherford model claim the electron should spiral inward. The reason for the inward direction is due to energy loss and the requirement of energy conservation.

But the motion of an electron is not controlled by a desire to satisfy energy conservation laws. The only thing that can make an electron move is force. And the only force that could cause it to spiral inwards is a retarding force against its direction of travel. Critics of the model make no suggestion of what that force could be. Nonetheless we know that an isolated hydrogen atom is unstable and this indicates its orbiting electron must either spiral in or away from the proton. Given that an inward spiral would involve releasing high-frequency radiation (as the spiral velocity increased), and we dont see this, it seems more likely that it spirals out. Furthermore an inward spiral would likely end with some sort of sub-atomic/nuclear particle interaction when the electron meets the proton, and we dont see this either. Hydrogen in a static field It happens that there is a way of making our above hydrogen model stable, and that is to place it within a static electric field. The field here will be net-neutral as shown below. On the left is a hydrogen atom consisting of a proton and electron. On the right is a positive and negative charge of equal but opposite value. These are far enough away to make the field approximately parallel.

When we set the electron in motion an interesting thing happens. See below. As before, the electron starts spiralling away from the proton. It then reaches a point where the orbit radius stops enlarging and becomes fixed in size. What has happened here? Due to the VDCL, as the electron moves right it feels increased repulsion from the negative charge and decreased attraction from the positive charge. Conversely as it moves left, attraction increases and repulsion decreases. The net result is a retarding force over the orbit period. When the electron was close to the proton it felt a promoting sideways force which pushed it away. But as it got further away, this force weakened and the retarding force from the distant charges took over. At some point these forces became equal and the orbit remained steady. If the electron instead starts outside this orbit it spirals inward and arrives at an (almost) identical orbit: This is very interesting because it shows that charged particles outside the atom can cause an electron to settle into a steady orbit at a fixed distance from a nucleus. Here the radius of orbit depends on the amount of charge in the exterior particles: increasing their charge reduces the radius and vice-versa.

So this is not really a fixed orbit like what Bohr might predict, but the above scenario does provide some clues as to how an orbit could become stable. A Simple Hydrogen Molecule One of the odd things about the development of atomic models is that fact that scientists always start by attempting to model the hydrogen atom. The reason this is odd is that it is well known by scientists that singular hydrogen atoms dont exist, other than very briefly. Ironically, the fact that the Bohr model predicts a stable hydrogen atom is proof that the model must be wrong! Instead hydrogen only becomes stable when joined to another atom; most commonly to another hydrogen atom in forming a H2 molecule. Using the VDCL and sideways force interactions, is it possible to arrange a pair of hydrogen atoms into stable molecular formation? How about this: Above we see two hydrogen atoms side-by-side with their electrons orbiting in mirrored but opposite directions. This arrangement solves several problems at once: 1. No emissions. When the electrons are orbiting in opposite directions, each radiates a sinusoidal signal that is the opposite of the other. The two signals cancel each other out, and will do so when observed from any direction.

In this way the molecule emits no net signal (no light). 2. No outward spiral As the electrons pass by each other they create a retarding sideways force. Much like the external charged particles described above, this prevents the orbits from becoming increasingly larger, i.e. from spiralling out. 3. A covalent bond This arrangement shows how a covalent (molecular) bond may be formed. When the electrons are at the far outer positions, the protons repel each other. Then when the electrons are near the mid section, each proton feels attraction toward the electrons. When these two forces balance the protons end up separated at a fixed distance, forming a covalent bond. Covalent bonds are often described with the term electron sharing which implies that each electron spends time orbiting both protons. Its important to note however that, while the electrons here do influence each others orbits, they only orbit their own proton, i.e. there is no sharing of electrons as such. 4. Fixed distance of orbit With the protons at a fixed distance, this in turn fixes the orbit radius of the electrons.

The electrons will settle into equilibrium when the promoting sideways force of the protons matches the retarding sideways force of the electrons. Here we see quantum mechanics come into play. A Hydrogen emission model The hydrogen molecule described above emits no radiation. This matches the normal situation we observe with unexcited hydrogen gas. What about the situation where excited hydrogen emits radiation, and does so at discrete frequencies? There may be a way of explaining this also. Above we see a hydrogen molecule in which the leftmost atom has lost its electron, perhaps as a result of being heated and colliding with another molecule. It might be supposed that the isolated proton should also break away at this point. But it will be held in place due to the fact that when the remaining electron is midway it creates an attractive force between the protons. The other thing this isolated proton can do is to keep the remaining electron from spiralling out. Due to the VDCL, the electron will feel a greater force moving away, than moving toward, the isolated proton. The result is a net retarding force which will keep the electron orbiting at a certain altitude. While in orbit, the electron will radiate a single frequency, possibly within the visible spectrum.

In the above diagram we note there are two degrees of freedom: the distance between the protons and the distance between the electron and its proton (the orbit altitude). We might suppose that there are only a certain number of possible distance solutions that satisfy an equilibrium situation involving a circular orbit and stable covalent bond. These solutions might correspond to the frequencies we observe in the hydrogen spectrum. Naturally the above is speculation and confirmation will require a proper analysis (to be done later). A Simple Helium Atom Next in the periodic table comes helium. Unlike hydrogen, helium is stable on its own and has no molecular form. Is it possible to construct a helium model using similar methods? How about this: Above is a nucleus containing two protons (charge=+2). It is surrounded on either side by electrons that are opposite each other and following the same circular orbit in clockwise fashion. Each electron emits a sine wave, but as the waves are 180 degrees out of phase, they cancel each others signal. As a result no radiation is emitted. This model is useful from the point of view of cancelling radiation. Unfortunately it has a number of problems: 1.

If this works for helium then it should also work for hydrogen, namely, it should be possible to place two electrons either side of a single proton. This would yield a negative hydrogen ion. Coulombs inverse square law tells us the opposing force between electrons will only be a quarter of the attractive force from the proton. So there is no problem force-wise. But we dont observe negative hydrogen ions. 2. The model only works if the electrons are exactly opposite each other and moving with equal speed. Any deviation from this will cause one electron to eventually move under and eject the other electron. 3. The most significant problem has to do with sideways forces. The promoting sideways force from the protons will cause electrons to spiral outward. This model intends to counteract that by making the electrons impose a retarding sideways force on each other. The electrons are further away but have the advantage of moving at twice the relative velocity. Will it be enough? Formula (2) above tells us that the velocities we are dealing with are around 1% of light speed. According to the earlier chapter on sideways force , at low velocities this force will be proportional to velocity and Coulombs force.

Let the velocity of an electron be 1, its charge be -1, and distance from the nucleus be 1. The nucleus (2 protons) then has a charge of +2 and a velocity of 0. The promoting sideways force from the proton then will be (setting c=1 and k=-1 for convenience): The electrons have a distance of 2 and relative velocity of 2. The retarding sideways force between them will be: As can be seen, the promoting sideways force is many times greater (4 times) than the retarding force. Hence the electrons will spiral out. If this model fails for helium, it would fail to an even greater degree with larger atoms, e.g. a lithium ion (Li+). In fact it also fails for hydrogen. Obviously we need a different model. How many neutrons does helium have? | Reference.com How many neutrons does helium have? A: Quick Answer The nucleus of helium contains two neutrons and two protons. Around the nucleus of helium there are two electrons in the 1s orbital. These have opposite spins. Full Answer Helium is produced in the sun, hence the name from the Greek sun god Helios.

Together with hydrogen, helium is one of the most common elements in the universe. Hydrogen is the simplest atom in its structure: it has only one proton, can have one electron and also one neutron. Helium is one step up in complexity, having two protons and two neutrons. Helium is an inert gas, which means that its electrons do not enter combinations with electrons of atoms of different elements or other helium atoms. Two helium atoms are attracted to each other only due to the force of electrical field that is created between them. However, at close distances the dominant reaction between two helium atoms is repulsion. The reason for the widely known voice-changing properties of helium is that sound travels approximately three times faster in helium than in the air. This raises the frequency at which the voice resonates and takes the pitch an octave and a half higher. It is not harmful to breathe helium in for these experiments as long as oxygen is inhaled as well. Q & A: Atom's Shells and Orbitals | Department of Physics | University of Illinois at Urbana-Champaign University of Illinois at Urbana-Champaign Q & A: Atom's Shells and Orbitals Learn more physics!

Most recent answer: 07/06/2016 Q: how do you determine how many shells an atom has and how many electron do those shells contain? - Anonymous (age 14) L.A. school district A: Good question! This is actually a very complicated question, and it's important to understand that orbitals and shells are basically just a good way to /think/ about what happens in the atom. Itï¿½s a bit truer to say that the electrons just buzz around in clouds around the nucleus, but this is definitely harder to think about. But hereï¿½s the basic idea for orbitals the way that youï¿½re probably learning about it in school: Every atom basically has an infinite number of shells. The thing is that almost all of those shells are empty (they donï¿½t have electrons in them). Electrons generally go into the orbital with the "lowest energy." The first orbital that fills up is called the 1S orbital. This one can hold 2 electrons. Once you get to the third electron, you have to put it in the next orbital. This one is called the 2S orbital.

(All "S" orbitals hold 2 electrons.) Next comes the 2P shell. "P" orbitals hold 6 electrons. After 2P is 3S, 3P, 4S, 3D, 4P, etc. ("D" orbitals hold 10 electrons each.) The order gets pretty complicated once you get this far, but you probably wonï¿½t need to know them much past 3P. If your teacher is trying to make this less confusing, they may just pay attention to the numbers, not the letters. In this case you can just add them up. Shell 1 (1S) holds 2 electrons. Shell 2 (2S & 2P) and Shell 3 (3S & 3P) each hold 8 electrons. After that, the next 2 hold 18 each, then the next 2 hold 32. So far, scientists havenï¿½t discovered any elements that use more orbitals than this. As for "how many shells the atom has," I already said that it has an infinite number of empty ones.

But the number of full ones is probably closer to what youï¿½re looking for. For example, lets look at Calcium. Calcium has 20 electrons. (You can see this by looking at the atomic number of Ca on a .) The first 2 go in shell 1, leaving 18 more. 8 more go into shell 2, then 8 in shell 3. Then there are 2 left. This isnï¿½t enough to fill the next shell, which would require 18 more, so we say that Calcium has 3 filled orbitals and 2 free electrons. If an atom is what is called an "ion," then that means it has a different number of electrons than youï¿½d expect. This is written by using little numbers with + or - after the atom's letters. Because electrons actually have a negative charge, a - sign means that the atom has more electrons than youï¿½d think. For example, letï¿½s take Calcium again. Like we said, Ca has 20 electrons. But Ca- has 21, and Ca2- has 22.

Going the opposite way, Ca+ has 19, and Ca2+ has 18. (Since 18 makes exactly 3 filled orbitals, Ca2+ is Calciumï¿½s most common ion.) Ok...going back to what I said about electrons going to the orbital with the "lowest energy." What if you were to give those electrons some more energy? Lets say we take Helium, which has 2 electrons. These 2 electrons will be just enough to fill Shell 1. If I give them some more energy, one of those electrons jumps up to Shell 2. So thereï¿½s 1 electron in Shell 1 and 1 in Shell 2. But as soon as it can, this electron will fall back to Shell 1 again. When it does, it will release energy as light. Hope this answers your question! -Tamara Kingston, Jamaica A: Youï¿½re on the right track. The two 2S electrons share exactly the same ï¿½orbitï¿½ but have opposite spins, so they arenï¿½t in the same exact quantum state.

The part thatï¿½s a little different from your picture is that the states these electrons are in arenï¿½t anything like classical orbits. Rather, the state is a wavelike cloud spread around the nucleus. For S-type states, the cloud has a perfectly spherical shape. For the 6 2P electrons, there are as you say 3 opposite-spin pairs of electrons, with the two electrons in each pair having the same ï¿½orbitï¿½, i.e. the same shape of cloud. The clouds in this case arenï¿½t spherical but rather have lobes. The three different 2P orbital states differ in that their lobes point different directions. Mike W. Clovis CA A: Yeah, that orbital stuff is complicated. I'm not sure why it gets introduced so soon in school. There would be a lot of much clearer things to learn. Just for a little start on orbitals, since something about them was probably assigned, you might look at these pictures of some different orbital shapes that the electron cloud can take: . Orbitals have stable shapes that don't change in time. Other states do change in time.

As for your wishes for science, perhaps providing governments with nuclear weapons was a step in that direction. If they are used, however, there will be collateral damage. Mike W. UK A: If each shell represented a spherically symmetrical cloud form ("S state"), you'd be right that each would hold only two electrons, given the two spin states. However, there are states that are not spherically symmetric. One way of expressing the P states has two lobes of cloud, with opposite sign of wave in each lobe. There are three independent ways of aligning those lobes, e.g. along x,y, or z axes. By symmetry these all have the same energy. For D states there are more lobes and five independent forms. In the conventional shell language, states that have almost the same energy, such as the lowest energy P states and the second-lowest-energy S states are lumped into the same shell, giving even more states per shell. Mike W.

Many-Electron Atoms - The Atomic Orbital Concept Many-Electron Atoms Table of Contour Values The Atomic Orbital Concept Since the nuclear charge is twice the electronic charge, the electrostatic energy of repulsion between the electrons will be the smallest of the three terms in equation (1) when the interactions are averaged over all possible positions of the electrons. We could obtain an approximation to the electronic energy of the He atom by neglecting this small term. This is a good idea for another reason. If we ignore the repulsion between the electrons, then physically we are supposing that neither electron "realizes" the other electron is present. The problem is thus identical to that of the hydrogen atom, a problem which can be solved exactly, except that the nuclear charge is now +2 rather than +1. The energy of each electron in the field of a nucleus of charge +2 is determined separately and the total energy is then simply the sum of the two energies. This is the first approximation to the energy. We will also obtain an approximation to the manner in which the electrons are distributed in space. With this latter knowledge, we can estimate the energy of repulsion between the electrons.

That is, with the knowledge of how the electrons are distributed relative to one another we can pick up the term in the potential energy we originally neglected, the term e2/r12, and calculate its contribution to the energy. This method of calculating the electronic energy of atoms reduces a many-electron problem to many one-electron problems. Each electron (moving in the attractive field of the nucleus) is treated independently of the others. In addition, since the problem is now a set of one-electron problems, we may carry over and use all of the results obtained for the hydrogen atom. We pointed out in our discussion of the hydrogen atom that the results we obtained could be applied to any one-electron system by setting Z equal to the appropriate value in all of the formulae. The one-electron energies are easily calculated and are given by . More important, the concept of atomic orbitals, the one-electron wave functions for the hydrogen atom, may be employed in the many-electron case. When each electron is considered in turn, its motion and distribution in space will again be determined by an atomic orbital.

The atomic orbitals will differ from the case of the hydrogen atom in that they will generally be more contracted In the previous chapter we pointed out that the average value of the distance between the nucleus and the electron, , decreased as the nuclear charge and hence the attractive force exerted by the nucleus was increased However, the orbitals will still be determined by the three quantum numbers n, l and m. Increasing Z contracts the orbital, but the symmetry of the problem is left unchanged, i.e., the attraction of the electron by the nucleus is still determined only by the distance between them and does not depend on the direction. The l and m dependences of the orbitals, which determine the directional properties of the orbital and of the electron density distribution, remain unchanged. Thus we may still refer to "hydrogen-like" 1s, 2s, 2p orbitals For example, the 1s orbital is the most stable orbital (most negative E value) for any value of Z, and we naturally assume that the most stable form of the helium atom will be obtained when both electrons are placed in the 1s orbital. This information, telling us in which atomic orbital the electrons have been placed, is called an electron configuration.

An abbreviated notation is used to denote the electron configuration For example, the lowest energy state of helium, in which two electrons are placed in the 1s orbital, is written as 1s2. (This is to be read as one-s-two and not as one-s-squared.) When one of the electrons is placed in an orbital of higher energy, an "excited" configuration is obtained. An example might be 1s12p1 which one electron is in the 1s orbital and one electron is in the 2p orbital. It should be emphasized that the concept of assigning an electronto an atomic orbital is a rigorous and exact concept only for the hydrogen atom; for the many-electron case it is an approximation. The atomic orbital approximation may be tested by applying it to the helium atom. We have seen that the energy of a single electron moving in the attractive field of a nucleus of charge +Ze is (2) The energy of the two electrons in the helium atom, each considered to be independent of the other, is simply for the 1s2 electron configuration. To this energy value must be added the energy of replusion between the two electrons.

Since both electrons have been placed in a 1s atomic orbital, we know that the charge distribution for each electron must be spherical and centred on the helium nucleus. The two charge distributions will be completely intermingled and we must calculate the energy of replusion between every small element of charge density of the one distribution with every small charge element of the second distribution. This calculation can be readily done by the methods of integral calculus and the value of the average energy replusion is found to be We label this energy Ec as it is a correction to our first approximation to the energy. Notice that in general Ec depends directly on the value of Z. This makes physical sense, for the greater the value of Z, the more contracted and superimposed are the two charge distributions, and the greater is the energy of replusion between them. Note as well that the correction Ec is indeed smaller than E1, an assumption we made in developing this method of approximating the electronic energy.

The estimate of the total electronic energy of helium atom is This total energy is called the electronic binding energy as it is the energy released when two initially free electrons are bound to the helium nucleus Recall that -K represents the binding energy of the most stable state of the hydrogen atom; thus the helium atom is five and one half times more stable than the hydrogen atom. This is not really a fair comparison because the value (11/2)K is the energy required to remove both electrons from the helium atom (a double ionization) It is more interesting to compare the energy required to remove a single electron from helium with the energy required to remove the single electron in hydrogen. The energy of the reaction (The energy required to ionize an atom will be denoted by the letter I.) and the total electronic energy of the lithium atom is predicted to be Thus it should require 214.4 ev to remove all three electrons from the lithium atom. We can also calculate the energy required to remove a single electron from lithium (the first ionization potential) When the predicted values for lithium are compared with the corresponding experimental values, they are found to be in serious error. The lithium atom is not as stable as the calculations would suggest.

Experimentally it requires 202.5 ev to remove all three electrons from lithium, and only 5.4 ev to remove one electron. Experimentally it requires less energy to ionize lithium than it does to ionize hydrogen, yet our calculation predicts an ionization energy one and one half times larger. The error in I is 300%! We should expect a realistic model to do better than this. More serious than this, however, is that the kind of calculation we are doing should never predict the system to be more stable than it actually is. The method should always predict an energy less negative than is actually observed. If this is not found to be the case, then it means that an incorrect assumption has been made or that some physical principle has been ignored. It is also clear that if we were to continue this scheme of placing each succeeding electron in the 1s orbital as we increased the nuclear charge by unity we would never predict the most striking property of the elements: the property of periodicity. We might recall at this point that there is a periodicity in the types of atomic orbitals found for the hydrogen atom. With every increase in n, all the preceding values of l are repeated, and a new l value is introduced.

If we could discover a physical reason for not placing all of the electrons in the1s orbital, but instead place succeeding electrons in the orbitals with higher n values, we could expect to obtain a periodicity in our predicted electronic structures for the atoms. This periodicity in electronic structure would then explain the observed periodicity in their properties. There must be another factor governing the behaviour of electrons and this factor must be one which determines the number of electrons that may be accommodated in a given orbital. To discover what this missing factor is and to find the physical basis for it, we must investigate further the magnetic properties of electrons. Electron Shells and the Bohr Model About Watch and Favorite Watch Watching this resources will notify you when proposed changes or new versions are created so you can keep track of improvements that have been made. Favorite Favoriting this resource allows you to save it in the “My Resources” tab of your account. There, you can easily access this resource later when you’re ready to customize it or assign it to your students. Electron Shells and the Bohr Model Niels Bohr proposed an early model of the atom as a central nucleus containing protons and neutrons being orbited by electrons in shells.

Learning Objective Construct an atom according to the Bohr model Key Points In the Bohr model of the atom, the nucleus contains the majority of the mass of the atom in its protons and neutrons. Orbiting the positively-charged core are the negatively charged electrons , which contribute little in terms of mass, but are electrically equivalent to the protons in the nucleus. In most cases, electrons fill the lower- energy orbitals first, followed by the next higher energy orbital until it is full, and so on until all electrons have been placed. Atoms tend to be most stable with a full outer shell (one which, after the first, contains 8 electrons), leading to what is commonly called the " octet rule ". The properties of an element are determined by its outermost electrons, or those in the highest energy orbital. Atoms that do not have full outer shells will tend to gain or lose electrons, resulting in a full outer shell and, therefore, stability. Terms Full Text Electron Shells and the Bohr Model As previously discussed, there is a connection between the number of protons in an element, the atomic number that distinguishes one element from another, and the number of electrons it has.

In all electrically-neutral atoms, the number of electrons is the same as the number of protons. Each element, when electrically neutral, has a number of electrons equal to its atomic number. An early model of the atom was developed in 1913 by Danish scientist Niels Bohr (1885–1962). The Bohr model shows the atom as a central nucleus containing protons and neutrons with the electrons in circular orbitals at specific distances from the nucleus . These orbits form electron shells or energy levels, which are a way of visualizing the number of electrons in the various shells. These energy levels are designated by a number and the symbol "n." For example, 1n represents the first energy level located closest to the nucleus. The Bohr model was developed by Niels Bohr in 1913. In this model, electrons exist within principal shells. An electron normally exists in the lowest energy shell available, which is the one closest to the nucleus. Energy from a photon of light can bump it up to a higher energy shell, but this situation is unstable and the electron quickly decays back to the ground state. In the process, a photon of light is released.

Electrons fill orbit shells in a consistent order. Under standard conditions, atoms fill the inner shells (closer to the nucleus) first, often resulting in a variable number of electrons in the outermost shell. The innermost shell has a maximum of two electrons, but the next two electron shells can each have a maximum of eight electrons. This is known as the octet rule which states that, with the exception of the innermost shell, atoms are more stable energetically when they have eight electrons in their valence shell, the outermost electron shell. Examples of some neutral atoms and their electron configurations are shown in . As shown, helium has a complete outer electron shell, with two electrons filling its first and only shell. Similarly, neon has a complete outer 2n shell containing eight electrons. In contrast, chlorine and sodium have seven and one electrons in their outer shells, respectively. Theoretically, they would be more energetically stable if they followed the octet rule and had eight. Bohr diagrams indicate how many electrons fill each principal shell. Group 18 elements (helium, neon, and argon are shown) have a full outer, or valence, shell. A full valence shell is the most stable electron configuration.

Elements in other groups have partially-filled valence shells and gain or lose electrons to achieve a stable electron configuration. An atom may gain or lose electrons to achieve a full valence shell, the most stable electron configuration. The periodic table is arranged in columns and rows based on the number of electrons and where these electrons are located, providing a tool to understand how electrons are distributed in the outer shell of an atom. As shown in , the group 18 atoms helium (He), neon (Ne), and argon (Ar) all have filled outer electron shells, making it unnecessary for them to gain or lose electrons to attain stability; they are highly stable as single atoms. Their non-reactivity has resulted in their being named the inert gases (or noble gases). In comparison, the group 1 elements, including hydrogen (H), lithium (Li), and sodium (Na), all have one electron in their outermost shells. This means that they can achieve a stable configuration and a filled outer shell by donating or losing an electron. As a result of losing a negatively-charged electron, they become positively-charged ions . When an atom loses an electron to become a positively-charged ion, this is indicated by a plus sign after the element symbol; for example, Na+.

Group 17 elements, including fluorine and chlorine, have seven electrons in their outermost shells; they tend to fill this shell by gaining an electron from other atoms, making them negatively-charged ions. When an atom gains an electron to become a negatively-charged ion this is indicated by a minus sign after the element symbol; for example, F-. Thus, the columns of the periodic table represent the potential shared state of these elements' outer electron shells that is responsible for their similar chemical characteristics. Electron Configuration | Wyzant Resources Resources / Lessons / Science / Chemistry / Electron Configuration Electron Configuration Electrons play a crucial role in chemical reactions and how compounds interact with each other. Remember, electrons are the negative particles in an atom that "orbit" the nucleus. Although we say they orbit the nucleus, we now know that they are actually in a random state of motion surrounding the nucleus rather than making circles around it, which is what an orbit implies. The best analogy to describe electron motion within an atom is how bees buzz around a beehive. They don't fly in complete circles around it, but they do hover and move around it in a seemingly random motion.

Electrons increase in elements as protons do, which is from left to right and from top to bottom on the periodic table. Therefore, the element with the fewest electrons would be in the top left-hand corner of the table and the element with the most electrons would be in the bottom right hand corner. The elements are arranged so that the increase from element to element is one electron. Therefore, in the first row, we see hydrogen and helium. This is because hydrogen has one electron and helium has two electrons, so we place them in ascending order. Electron Orbitals We categorize electrons according to what orbital level in which they reside. The four orbitals are s, p, d, and f. They are classified by divisions on the periodic table, as follows: The first orbital is the s orbital. It has room to hold two electrons. The electrons have opposite spins, so it makes sense that they are paired together. The s orbital is a sphere, with the x, y, and z axes passing through it, like this: This means that the two electrons can occupy any of the space seen in this sphere, and they sort of "hover" around in the given space. The next orbital is the p orbital.

It can hold up to six electrons, therefore it has three sub-orbitals (each can hold two electrons). The spins on electrons are still opposite, this time split into three and three (since the first orbital only held two electrons, we said the spins were opposite. Now that this orbital can hold six electrons, three spin one way and three spin the opposite way). The p orbital is not sphere shaped, however it does have six lobes that are shaped like balloons. Two lobes are on the x axis, two are on the y axis, and two are on the z axis. These three separations are considered sub-orbitals and combine to make up the entire p orbital. The nucleus of the atom is located where these three axes meet. The p orbital looks like this: The next orbital is the d orbital. It can hold up to 10 electrons, therefore it has five sub-orbitals (each can hold two electrons). The spins of the electrons are opposite, so five are spinning one way and the other five are spinning the opposite way. The d orbital is not sphere shaped; it looks more like the p orbital, except there are more lobes that cannot be shown all at once.

We showed the entire p orbital (all three of the sub-orbitals) in one diagram, because there were two lobes on each axis. However, we need to show the five different sub-orbitals of the d orbital in order to fully explain where the lobes are located, and how they are shaped. We will show you four views, with labels on all of the axes. The first view is of the lobes that lie on the XY plane, shown in aqua here. The second view is a three dimensional view of lobes on the Z axis that rotate 360 degrees around the axis. There are two lobes, one in the top hemisphere and one in the bottom, and a tube-shaped area that circles the Z axis and intersects the X and Y axes. It's shown here in orange. The third view is of the lobes on the ZY plane, with the X axis running perpendicular to it. It's shown here in green. The last view is of the lobes lying on the ZX plane, and is shown here in pink. If all of these layers were put together, we would see a sort of star-burst image, with a tube encircling the middle.

The final orbital is the f orbital, and scientists are not completely sure of the shape of its orbital. However, they do have seemingly accurate predictions of where electrons will fall. We will show you the following probabilities of where electrons lie: We showed you two probabilities of where the f orbitals lie; however, the first image (in blue) is shown on the Z axis. It is actually repeated on the X axis and again on the Y axis. The second image (in orange) is shown in the XYZ dimensions; however, it is repeated three more times for a total of four positions using this shape and lobe configuration. We say that these are probable locations because scientists cannot actually track and determine the exact location of electrons. However, through research and abilities to track electrons in other orbitals, scientists can say that the likely location of f-level electrons is in one of these locations. Diagonal Rule, or Madelung's Rule In chemistry, the Diagonal Rule (also known as Madelung's Rule) is a guideline explaining the order in which electrons fill the orbital levels. The 1s2 orbital is always filled first, and it can contain 2 electrons.

Then the 2s2 level is filled, which can also hold 2 electrons. After that, electrons begin to fill the 2p6 orbital, and so on. The diagonal rule provides a rule stating the exact order in which these orbitals are filled, and looks like this: As you can see, the red arrows indicate the filling of orbital levels. Starting at the top, the first red arrow crosses the 1s2 orbital. If you follow these arrows down the list, you can easily determine the order that electrons fill the orbital levels. There is an exception to this rule when filling the orbitals of heavier electrons. For example, when filling the 5s2 orbitals, the rule says that 5s2 will fill, and then 4d10 will fill. However, when filling these orbitals for certain metals, only one electron will fill the 5s2 orbital, and the next electron will jump into the 4d10 orbital. This can be predicted, but cannot be exactly determined until it is observed.

The same is true for the 6s2 orbital-for certain heavy metals, the 6s2 will only contain one electron, and the other electrons will jump to the 5d10 orbital. Electron Notation Following the diagonal rule, there is an easy way to write electron configuration. We are simply going to use the orbital names we learned from the diagonal rule (1s2, 2s2 and so on). However, we are only going to write the number of electrons that the atom actually contains. For example, hydrogen has one electron, which would fall in the 1s orbital. Thus, the electron configuration for hydrogen is 1s1. We write the superscript as 1 because there is one electron. Helium, the next element, contains two electrons. They both fill the 1s orbital, so the electron configuration for helium is 1s2. Here again, we write the superscript as 2 because there are two electrons. Electron configuration moves across and down the periodic table. You might have noticed that we first put one electron in the 1s orbital (with hydrogen), and then we put two electrons in the 1s orbital (with helium).

Continuing this trend, we would next have 3 electrons with lithium. We would place two of them in the 1s orbital, and one of them in the 2s orbital, so the electron configuration would be 1s2 2s1. However, we can also write this using the configuration of helium, because it is a noble gas. Noble gases are stable elements, so we can use their configurations in determining other configurations. So, instead of writing 1s2 2s1, we would write [He] 2s1. This means that lithium contains the same configuration as helium, and then has one more electron in the 2s orbital. Notice that we use brackets to encase the previous noble gas, and then we continue writing the configuration as we normally would. This might not seem like a big deal, or a shortcut right now, but once you get pretty far down the periodic table, this will save you a lot of time and energy. We'll show you one short example of this. Let's say we need to determine the electron configuration for Ba, Barium.

Counting across the table, we would come up with the following configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d 10 4p6 5s2 4d10 5p6 6s2 Instead of writing all of this out, we could simply find the previous noble gas, which is Xe. Therefore, we can write [Xe] and then figure out the rest of the configuration. We look and see that Ba is in the 6th row, so we know that we're going to start with 6s2. We can look and see that Ba is the second element in that row, so it has two electrons to go in the 6s orbital. Thus, we can conclude that the final electron configuration for Ba is [Xe] 6s2. Electron Spin Every electron placed in an orbital has a feature we refer to as "spin." We've already talked about electrons having been thought to have a specific "orbit," and then later discovered to hover in the places listed above.

Well, electrons don't literally spin, but their movement sort of looks like someone somersaulting, in a very fast, random state. This is called spin. Each electron can either have a +1/2 spin or a -1/2 spin, which indicates its direction of motion. There is never a 0 spin. When filling orbitals, electrons spin in pairs, one + and one -. The electrons with an up (+) spin fill first, and the electrons with a down (-) spin fill second. It would look like this: and so on. Because the s orbital can hold two electrons, we draw one box in which two electrons (represented here with arrows) can fit. Since the p orbital can hold 6 electrons, we draw 3 boxes that will each hold 2 electrons. This will continue with the d orbital (10 electrons fit in 5 boxes) and the f orbital (14 electrons fit in 7 boxes). These boxes fill in a certain order-all of the boxes in one column will fill with up spin electrons first, and then down spin.

So, for example, if the element has a configuration of 1s2 2s2 2p3, it would look like this: As you can see, we filled the boxes with up arrows (electrons) first. If we had more electrons, we would go back and add them in to the second column, as down spin electrons. Sign up for free to access more Science resources like . Wyzant Resources features blogs, videos, lessons, and more about Science and over 250 other subjects. Stop struggling and start learning today with thousands of free resources! Electron Pairing Electron Pairing After completing this section you will be able to do the following: Explain what paired and unpaired electrons are. Explain how paired electrons affect the magnetic properties of a material. Questions What are paired electrons? How do paired electrons affect the magnetic properties of a material? Since all matter is made up of atoms and all atoms have electrons that are in motion, do all atoms have magnetic fields? The answer to this question is yes and no.

All the electrons do produce a magnetic field as they spin and orbit the nucleus; however, in some atoms, two electrons spinning and orbiting in opposite directions pair up and the net magnetic moment of the atom is zero. Remember that the direction of spin and orbit of the electron determines the direction of the magnetic field. Electron pairing occurs commonly in the atoms of most materials. In the experiment you observed a helium atom showing two electrons spinning and orbiting around the protons and neutrons of the nucleus. The two electrons are paired, meaning that they spin and orbit in opposite directions. Since the magnetic fields produced by the motion of the electrons are in opposite directions, they add up to zero. The overall magnetic field strength of atoms with all paired electrons is zero. In general, materials that have all paired electrons in the atoms and thus have no net magnetic moment are called diamagnetic materials; yet, there are some exceptions. When placed in the magnetic field of a magnet, diamagnetic materials will produce a slight magnetic field that opposes the main magnetic field. Both ends of a bar magnet will repel a diamagnetic material.

If a diamagnetic material is placed in a strong external magnetic field, the magnetic field strength inside the material will be less than the magnetic field strength in the air surrounding the material. The slight decrease in the field strength is the result of realignment in the orbit motion of the electrons. Diamagnetic materials include zinc, gold, mercury, and bismuth. Another key concept in magnetism is that diamagnetic materials will oppose an applied magnetic field. Both ends of a magnet will repel diamagnetic materials. Are all materials that have unpaired electrons magnetic? Most materials with one or more unpaired electrons are at least slightly magnetic. Materials with a small attraction to a magnet are called paramagnetic materials, and those with a strong attraction are called ferromagnetic materials. Aluminum, platinum, and manganese are some paramagnetic materials. Iron, cobalt, and nickel are examples of ferromagnetic materials. ATOMS ATOMS All matter is made from particles called atoms. The Ancient Greeks described atoms as the smallest particles that make up everything. We now know that atoms are built from even smaller particles. Atoms link together and make MOLECULES . HOW BIG IS AN ATOM?

The radius of a typical atom is one tenth of a billionth of a meter. A string of atoms one meter (about 3 ft) long contains an atom for every person in the world. A cube of sugar contains as many atoms as there are stars in the Universe. The biggest atom (cesium) is nine times the diameter of the smallest atom (helium). WHAT’S INSIDE AN ATOM? The particles that make up atoms are electrons, protons, and neutrons. Their position in an atom is the ATOMIC STRUCTURE . Electrons are a kind of lepton. Protons and neutrons are made up of three quarks each. Quarks and leptons are fundamental particles—the smallest particles in the Universe. HOW CAN ATOMS BE SPLIT? Protons and neutrons are held together in the nucleus at the center of the atom by a strong force. But this force can be overcome by striking the nucleus with a neutron, a proton, or another particle. The nucleus may split and form new atoms. Atoms are split in this way inside nuclear reactors and during nuclear explosions. ATOMIC BOMB In a nuclear explosion, the atom’s nucleus (center) is split.

Neutrons are released, and hit other nuclei, setting off a chain reaction. The result is a big release of energy in a blast of heat, light, and radiation. ATOMIC STRUCTURE Most of an atom is empty space. Protons and neutrons occupy the nucleus at the center of the atom. Electrons orbit the nucleus like planets around a star. They are grouped in layers called shells. WHAT MAKES ATOMS STICK TOGETHER? Electrons carry a negative electric charge, and protons carry a positive charge. The attraction between them holds electrons in orbits. When atoms come together, they share electrons in their outer shells to form chemical bonds. This model of a carbon atom is split in half to show the inside. The nucleus of a carbon atom contains six neutrons and six protons. Six electrons orbit the nucleus in two shells. MOLECULES Different atoms bonded (stuck) together in particular arrangements are called molecules. Water molecules, for example, have two hydrogen atoms bonded to one oxygen atom. WHAT SHAPES ARE MOLECULES? The simplest molecules consist of just two atoms, and are shaped like dumbbells.

But atoms can link together to make molecules of almost any shape imaginable—pyramids, chains, rings, spirals, balls, or tubes. FIND OUT MORE Many-electron atoms Many-electron atoms The hydrogen atom is the only atom for which exact solutions of the Schrödinger equation exist. For any atom that contains two or more electrons, no solution has yet been discovered (so no solution for the helium atom exists!) and we need to introduce approximation schemes. Let us consider the helium atom. The nucleus has a charge of , and if we place the nucleus at the origin, there will be an electron at a position with spin and an electron at position and spin . As usual, we consider the nucleus to be fixed. The classical energy is then Note that is not simply a sum of terms for electron 1 and electron 2, . Therefore, it is not possible to write the wave function as a simple product of the form , nor is it even possible to use the special form we introduced for identical particles because these simple products are not correct solutions to the Schrödinger equation.

This means that the wave function depends on the full set of 6 coordinates or if spherical coordinates are used, and 2 spin coordinates , , and that this dependence is not simple! In fact, as the number of electrons increases, the number of variables on which depends increases as well. For an atom with electrons, the wave function depends on coordinates! Thus, it is clear that the wave function for a many-electron atom is a very unwieldy object! As a side bar, we note that the 1998 Nobel prize in chemistry was awarded to Walter Kohn for the development of an extremely elegant theory of electronic structure known as density functional theory. In this theory, it is shown that the wave function , which depends on coordinates, can be replaced by a much simpler object called the electron density denoted . This object depends on only three variables for a system of any number of electrons. In density functional theory, it is shown that any physical quantity can be computed from this electron density . Consider now an imaginary form of helium in which the two electrons do not interact. For this simplfied case, the energy is simply

Question: How many electrons orbit the nucleus in one atom of Helium?
Answer: TWO